University

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University POLITEHNICA of Timisoara Faculty of
Industrial Chemistry and Environmental
Engineering
GENERAL CHEMISTRY
Assoc.Prof. dr.eng. Andrea Kellenberger
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GENERAL CHEMISTRY COURSE
Lecture 2h / week Laboratory 2h /
week Evaluation form Examination Nr. of credits
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Chapter 1
INTRODUCTION
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Einsteins equation
E mc2
E energy, in J m mass, in kg c light speed,
in m s-1
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Substance is what all things consists of. One of
the most important properties of the substance is
the mass. Other characteristics of the substances
are homogeneity and constant composition. Homogene
ity same characteristics in all the
volume. Constant composition in any given
portion of a substance there are the same
particles which interact in the same
way. Examples of substances water, sugar,
oxygen, sodium chloride, copper, hydrochloric
acid, sodium hydroxide.
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The great majority of things consist of a
mixture of substances. The air, for example, is
not a substance, because it can be separated
through distillation in oxygen, nitrogen, argon
and other gases. The petrol is not a substance as
well through distillation it can be separated
in hydrocarbons. Chemistry is the science of
substances. Chemistry studies the structure,
properties and changes of substances.
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International System of Units (
SI units) (from French Le
Systeme International dUnites). Making
observations is a key part of the scientific
process. Sometimes observations are qualitative,
for example the substance is a green gas, and
sometimes they are quantitative (the mass is 10
grams, the temperature is 25ºC or the pressure is
1015 mbars). A quantitative observation is called
a measurement. A measurement always consists of
two parts a number and a unit.
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Table 1. SI Fundamental Units
Quantity Unit Symbol
Length meter m
Mass kilogram kg
Time second s
Thermodynamic temperature Kelvin K
Electric current Ampere A
Luminous intensity candela cd
Amount of substance mole mol
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Table 2. SI Derived Units
Quantity Quantity equation Name of the unit Symbol
Area S l2 l length square meter m2
Volume V l3 cubic meter m3
Speed v l / t t time meter per second m s-1
Acceleration a v / t meter per second squared m s-2
Molar concentration cM n / V n- amount of substance V-volume mole per cubic meter mol m-3
Force F ma m mass a acceleration kilogram meter per second squared (Newton) kg m s-2( N )
Pressure P F / S S area kilogram meter per square second per square meter kg m s-2 m-2 N m-2 Pa
Density ? m / V kilogram per cubic meter kg m-3
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Examples - for volume (liter, L 1L 1dm3), for
pressure (bar 1 bar 105 Pa) -
for temperature, there are two non SI scales
Celsius (ºC) and Fahrenheit (ºF). Temperature
conversions can be made using the equations shown
below Kelvin from Celsius
T (K) t (ºC) 273.15 Celsius from Fahrenheit
t (ºC ) 5/9 t(ºF) 32
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CHAPTER 2 ATOMIC STRUCTURE OF THE SUBSTANCES
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One of the oldest scientific concepts is
that all matter can be broken down until finally
the smallest possible particles are reached
these particles cannot be further subdivided. The
Greek philosopher Democritus (460 370 B.C.)
considered these particles to be in a constant
motion, but able to fit together into stable
combination. The characteristics of substances
resulted from the different size, shape and
arrangements of the particles, named atoms. In
Greek, atom means indivisible.
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• Modern atomic theory was developed by John Dalton
  (1776 1844), based on quantitative data, not
  only qualitative observations or speculations.
  Two natural laws serve as the basics of Daltons
  atomic theory
• Law of conservation of mass
• Law of definite composition (proportions)

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a. Law of conservation of mass The total mass
of materials present after a chemical reaction
is the same as before the reaction. Example
the reaction between Mg and O2. 0.24g Mg react
with 0.16g oxygen. After the reaction, 0.40g MgO
(magnezium oxide) were obtained. b. Law of
definite composition (proportions) All
samples of a compound have the same composition,
the same proportions by mass of the constituent
elements.
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To see how the law of constant composition
works, lets consider the compound FeS (iron
sulfide). A sample of 10g FeS contains 6.36g Fe
and 3.64g S. That means
and
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Another sample of 25g FeS contains 15.91g Fe
and 9.09g S. The composition of the second sample
is the same
and
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• J. Dalton was aware of these observations and he
  offered an explanation for them. This is known as
  Daltons atomic theory. The main ideas of this
  theory can be stated as follows
• Chemical elements are made of small particles
  called atoms.
• All atoms of a given element are identical.
• The atoms of a given element are different from
  those of any other element.

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• Atoms of one element can combine with atoms of
  other elements to form compounds. A given
  compound always has the same relative number and
  type of atoms.
• Atoms are indivisible in chemical processes.
  That is, atoms are not created or destroyed in
  chemical reactions. A chemical reaction simply
  changes the way the atoms are grouped together.

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In order to apply Daltons theory in predicting
new phenomena, it was necessary to assign
characteristic masses to atoms. These masses
became known as atomic weights. Since they are
very small, it is impossible to isolate and weigh
individual atoms. Dalton tried to establish
relative atomic weights. If an atom of hydrogen,
for example, is taken to have the mass of 1 unit,
the mass of oxygen atom is 16.
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2.1. Atom structure
At the end of the 19th century, the English
physicist J. J. Thomson showed that the atoms of
any element emit tiny negative particles. Thus,
he concluded that all types of atoms must contain
these negative particles called electrons.
Although atoms contain negative particles, the
whole atom is not negatively or positively
charged. So, Thomson concluded that the atom must
also contain positive particles that balance the
negative charge given by the electrons. He
imagined the plum pudding model of the atom, in
which electrons are scattered like plums into the
uniform pudding of positive charges.
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Rutherfords experiment Rutherford measured
the deviation of alpha particles (helium ions
with a positive charge) directed normally onto a
sheet of very thin gold foil. Assuming the plum
pudding model, the alpha particles should all
have been deviated by, at most, a few degrees.
The results of the experiment were very different
from those Rutherford anticipated. Most of the a
particles passed straight through the foil, some
of them were deflected at large angles and some
were even reflected backward.
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Rutherfords experiment
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• Conclusions
• the plum pudding model for the atom could not be
  correct
• the atom must contain a very small (compared
  with the size of the atom) positive charge which
  causes the large deflections of the a particles
• the atom is mostly empty space because most of
  the a particles past directly through the foil.

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These results could be explained only in terms
of nuclear atom an atom with a dense center of
positive charge (nucleus) around which tiny
electrons moved in a space otherwise empty. He
concluded that the nucleus has a positive charge
to balance the negative charge of the electrons
and that it must be small and dense. This picture
of the atom was the planetary model of the atom.
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Planetary model of the atom
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In 1919 Rutherford discovered that the nucleus
contains positive particles named protons.
Furthermore, in 1932, James Chadwick discovered
that nucleus contains also neutral particles
neutrons. Protons and neutrons are known as
nucleons. Electrons, protons and neutrons are
fundamental particles of the matter. There are
more than 30 other fundamental particles.
Properties of the main fundamental particles are
given in the next table.
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Properties of the main fundamental particles
Electric charge Electric charge Mass Mass Symbol
C Relative charge kg amu Symbol
Proton 1.60210-19 1 1.6726210-27 1.0073 p p
Neutron 0 0 1.6749310-27 1.0087 n no
Electron -1.60210-19 -1 9.1093910-31 0.0005486 e e-
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• The attraction force between the positive
  charges (protons) and the negative ones
  (electrons) keeps the atom together. In this
  image the nucleus is like a sun and electrons
  like planets. Several problems arise with this
  concept the electrons might be expected to slow
  down gradually and fall on the nucleus?
• To explain why this did not occur, Niels Bohr
  (1913) postulated
• The electrons can move around the nucleus only
  on certain orbits (allowed orbits).

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• The electrons can gain or lose energy only by
  jumping from one allowed orbit to another. When
  an electron moves towards the nucleus energy is
  radiated and if it moves away from the nucleus
  energy is absorbed.
• For an electron to remain in its orbit the
  electrostatic attraction force between the
  electron and the nucleus must equal to the
  centrifugal force which tends to throw the
  electron out of its orbit.
• N. Bohr admitted that the orbits of the electrons
  are circular.

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Bohrs model of the atom
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A. Sommerfeld (1916), based on the atomic
spectra of hydrogen, suggested that the
permissive orbits of the electrons may be
elliptic.
Bohr Sommerfeld model of the atom
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Atomic number, mass number and chemical element
The number of protons in an atom is called the
atomic number Z. In an atom, which must be
electrically neutral, the number of electrons are
also equal to Z. The total number of protons and
neutrons in an atom is the mass number A. Thus,
the number of neutrons is A-Z. The three
subatomic particles considered, the electron,
proton and neutron, are the only ones involved in
chemical phenomena. A study of matter at its most
fundamental level must consider a lot of
additional subatomic particles.
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All atoms with the same number of protons
signify a chemical element. Each element has a
name and a distinctive symbol. Chemical symbols
are one or two letter abbreviations of the
elements name (usually the Latin name). The first
letter, but never the second is capitalized. For
example
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Hydrogenium H Aurum - Au
Nitrogenium N Cuprum - Cu
Carbonum C Silicium - Si
Oxygenium O Ferrum Fe
Phosphorus P Tellurium - Te
Sulphur S Natrium - Na
Fluorum - F Aluminium - Al
Hydrargirum Hg Strontium - Sr
Stibium Sb Protactinium - Pa
Platinum - Pt Plutonium Pu
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To represent the composition of any particular
atom, we need to specify its number of protons,
neutrons and electrons. We can do this with the
symbol
Atoms that have the same atomic number Z, but
different mass numbers A are called isotopes.

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Hydrogen has as well another two isotopes
Natural abundance of hydrogen isotopes is
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Abundance of the elements
What is the most abundant element? This simple
question does not have a simple answer. If we
consider the entire Universe, hydrogen accounts
for about 90 of all the atoms and 75 of the
mass, and helium accounts for most of the rest.
If we consider only the elements present on
Earth, iron is probably the most abundant
element. However, most of the iron is in Earths
core. The currently accessible elements are those
present in Earths atmosphere, oceans and solid
continental crust up to 16 km depth. The relative
abundance in these parts of the Earth are called
Clark parameters.
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Nr. crt. Element Clark Parameter Nr. crt. Element Clark Parameter
1 Oxygen 49,4 16 Samarium 5?10-4
2 Silicon 25,75 17 Gadolinium 5?10-4
3 Aluminum 7,51 18 Dysprosium 5?10-4
4 Iron 4,7 19 Ytterbium 5?10-4
5 Calcium 3,39 20 Erbium 4?10-4
6 Sodium 2,64 21 Argon 3,6?10-4
7 Potassium 2,40 22 Praseodymium 3,5?10-4
8 Magnesium 1,94 23 Lutetium 1?10-4
9 Hydrogen 0,88 24 Germanium 1?10-4
10 Titanium 0,58 25 Selenium 8?10-5
11 Chlorine 0,19 26 Cesium 7?10-5
12 Phosphor 0,12 27 Terbium 7?10-5
13 Carbon 0,087 28 Holmium 7?10-5
14 Manganese 0,085 29 Thulium 7?10-5
15 Sulfur 0,048 30 Niobium 4?10-5
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Not all the known elements exist in Earths
crust. There are only 88 natural elements. The
rest of known elements can be produced only
artificially by nuclear processes. Moreover, most
of the elements do not occur free in nature, that
is, as uncombined element. Only about 20 of them
do. The remaining elements occur in chemical
combinations with other elements. We can see in
the last table that oxygen is the most abundant
element in the Earths crust (49.4).
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There are 3 natural isotopes of oxygen
Large amount of oxygen exists in water and rocks
as well in free state like molecular oxygen (O2)
and ozone (O3). Molecular oxygen (O2) and ozone
(O3) are allotropes of the element oxygen. The
second element in Clarks table is silicon (Si
25.75), but silicon occurs only in chemical
combinations.
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Atomic mass unit
Because the fundamental particles and the atoms
are very tiny it is difficult to operate with
the small values of their masses. This is the
reason why the atomic mass unit (amu) was
introduced.
1 amu 1.66 10-27 kg
Generally, atomic masses of the elements are
fractional number because the natural elements
are a mixture of two or more isotopes.
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For example, magnesium has 3 stable isotopes
(78,70 ), exact atomic mass 23,98504
(10,13 ), exact atomic mass 24,98384
(11,17 ), exact atomic mass 25,98259
Knowing the abundance of the stable isotopes one
may calculate the atomic mass of magnesium
AMg0.787 x 23.98504 0.1013 x 24.98384 0.1117
x 25.98259 24.30934
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Electronic configuration of the atoms
Louis de Broglie (France) and Werner Schrödinger
(Austria) in the mid 1920s, suggested that like a
light, the electron has both a wave and particle
properties. When Schrödinger carried out a
mathematical analysis based on this idea, he
obtained a new model for the atom wave model.
In this model the electron has not a well
defined orbit. The motion of the electron seems
to be rather a vibration. The three-dimensional
region of space around the nucleus in which we
can find the electron is called orbital. In fact,
it is a region of probability where the electron
is likely to be found.
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Let us consider a multi electronic atom. We can
assume that each electron has a specific mean
path from nucleus. The electrons having a similar
mean path form a main energetic level or main
electronic shell, characterized by the principal
quantum number n. The main shells are denoted by
letters.
Energetic level K L M N O P Q
Principal quantum number n n 1 n 2 n 3 n 4 n 5 n 6 n 7
The first main shell is the nearest level to the
nucleus and it has a minimum energy.
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K shell (n 1) consists of one s sublevel,
containing one orbital with a spherical symmetry
named s orbital
s orbital
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L shell (n 2) has 2 sublevels one s sublevel
containing one spherical shaped s orbital and one
p sublevel containing 3 p orbitals. Each p
orbital consists in two lobes distributed along
one of the three rectangular axes through the
nucleus
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In order to characterize the shape of the
orbital the orbital quantum number or azimuthal
number l has been introduced. For s orbital l 0
and for p orbital l 1. All orbitals having
the same l value form a subshell or
sublevel. The orientation of the orbitals is
given by the magnetic quantum number m. It may be
0 1 2 l. For example, if l 1 provided
that m -1 0 1, that is there are 3 p
orbitals px, py and pz.
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For the third electronic level M (n 3), the
values of the orbital and magnetic numbers are
the following
n 3 l 0 m 0
n 3 l 1 m -1 0 1
n 3 l 2 m -2 -1 0 1 2
Beside s and p orbital, on the third electronic
level there are 5 orbitals characterized by
orbital number l 2, named d orbital. There are
5 different d orbitals
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The d orbitals
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Furthermore, for the 4th electronic level N, the
quantum numbers are
n 4 l 0 m 0
n 4 l 1 m -1 0 1
n 4 l 2 m -2 -1 0 1 2
n 4 l 3 m -3 -2 -1 0 1 2 3
In this case, the orbitals with l 3 are called
f orbitals. According to the values for m, there
are 7 f orbitals
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The f orbitals
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The energy of the levels increases as the values
of the n increase. An orbital can be empty or it
can contain one or maximum two electrons. If two
electrons occupy the same orbital, they must have
opposite spins, associated with spin quantum
number s, which may be 1/2.
The order of filling orbitals
The electron configuration is the arrangement
of electrons on shells, subshells and orbitals.
Electrons fill low energy orbitals, closer to the
nucleus, before they fill higher energy ones.
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The order of energy levels is not identical to
the principal quantum number, due to the
interaction between electrons and nucleus. The
attractive force of the nucleus for a given
electron increases as the nuclear charge
increases. Therefore, the orbital energy
should decrease with increasing the atomic
number. The energy for the principal quantum
levels are given in the next figure.
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The order in which electrons occupy orbitals
is 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d
6p 7s 5f 6d 7p In order to establish the
electronic configuration of an atom a lot of
algorithms have been proposed. One of these
algorithms is known as a minimum (n l) rule.
According to this rule, the electronic levels and
sublevels will be filled in the increasing order
of the sum of the principal quantum number and
the orbital one. For the same sum n l, the low
energy corresponds to the orbital with the lower
n.
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Minimum n l rule can be illustrated by the
following table
n 1 2 2 3 3 3 4 4 4 4 5 5 5 5 5 6 6 6 6 6 6
l 0 0 1 0 1 2 0 1 2 3 0 1 2 3 4 0 1 2 3 4 5
s s p s p d s p d f s p d f g s p d f g h
n??l 1 2 3 3 4 5 4 5 6 7 5 6 7 8 9 6 7 8 9 10 11
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A suggestive image of the n l rule has been
given by Goldanskys chessboard
The Goldanskys chessboard
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Paulis exclusion principle Two electrons in
an atom cannot have all four quantum numbers
alike. If two electrons exist in the same
orbital, that is they have identical principal
number, orbital and magnetic ones, these
electrons must have opposite spins (different
spin numbers). Maximum number of electrons in an
electronic shell with principal number n is 2.
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Hunds rule When orbital of identical energy
are available, electrons occupy these singly
rather than in pairs. As a result, an atom tends
to have as many unpaired electrons as
possible. In its ground state, hydrogen has its
electron in the 1s orbital. This is commonly
represented in two ways the electron
configuration or orbital diagram (box diagram). A
small arrow indicates the electron.
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1s H 1s1 ?
Electron configuration orbital (box) diagram
The next element is helium (He). It has two
protons in the nucleus and so it has two
electrons. Both electrons are placed in the 1s
orbital, but they have opposite spins.
1s He 1s2 ??
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Lithium (Z 3) has three electrons, two of
which go into the 1s orbital. That is, the 1s
orbital is full. The third electron must occupy
an orbital with n 2, in this case 2s orbital.
1s 2s Li 1s22s1 ?? ?
The next element, beryllium (Z 4), has 4
electrons which occupy the 1s and 2s orbitals
with opposite spins
1s 2s Be 1s22s2 ??
??
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Boron (Z5) has five electrons, four of which
occupy the 1s and 2s orbitals. The fifth electron
goes into the second type of orbitals with
principal quantum number 2, one of the 2p
orbitals
1s 2s 2p B 1s22s22p1
?? ?? ?
Because all the 2p orbitals have the same
energy, it does not matter which 2p orbital the
electron occupies. Carbon (Z6), the next
element, has six electrons two in 1s orbital,
two in 2s orbital and two in 2p orbital. The last
electrons occupy different 2p orbitals
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1s 2s
2p C 1s22s22p2 ?? ?? ? ?
The configuration of nitrogen (Z 7) is
1s 2s 2p N 1s22s22p3 ??
?? ? ? ?
Oxygen (Z 8) has four electrons in the 2p
orbitals, one of the 2p orbitals is occupied by a
pair of electrons with opposite spins

1s 2s 2p O 1s22s22p4
?? ?? ?? ? ?
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The electron configuration for fluorine (Z 9)
and neon (Z 10) are
1s 2s 2p F 1s22s22p5 ?? ??
?? ?? ? Ne 1s22s22p6 ?? ?? ??
?? ??
With neon, the K (n 1) and L (n 2) shells
are filled. For sodium (Z 11), the eleventh
electron fill the first orbital of the M (n 3)
shell, that is 3s orbital. Thus, the electron
configuration for sodium will be 1s22s22p63s1. To
avoid writing the inner-level electrons the
configuration of neon 1s22s22p6 is abbreviated
with Ne.
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1s
2s 2p 3s Na
Ne3s1 ?? ?? ?? ?? ?? ?
We have to note that the configuration of an
element differs from the previous only by an
electron named differentiating electron.
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Periodic table of the elements
The chemical and physical properties are
determined by the number and arrangement of the
electrons, that is by the atomic number. If the
elements are arranged in groups, each group
having a characteristic electronic structure,
then elements should show similarities in
chemical and physical properties. A
classification scheme of the elements, similar to
that used today, was discovered independently by
Dmitri Mendeleev and Luther Meyer in 1869. Their
classifications were based on an early version of
the periodic law
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If the elements are arranged in order of
increasing atomic mass, certain sets of
properties are found to reappear
periodically. The arrangement of the elements
based on the periodic law is called periodic
table. In Mendeleevs periodic table the elements
were arranged in 12 horizontal rows and 8
vertical columns or groups. The eight groups were
further divided in 16 sub-groups. Mendeleevs
periodic table was a short form. The modern
periodic table is a long form.
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The horizontal rows of the table are called
periods. The first period of the table consists
of only two elements hydrogen (H) and helium
(He). The second and third periods have eight
elements each from lithium (Li) to neon (Ne)
and from sodium (Na) to argon (Ar). The fourth
and fifth periods comprise 18 elements each,
ranging from potassium (K) to krypton (Kr) and
from rubidium (Rb) to xenon (Xe).
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The sixth period is a long one with 32 members.
From this period 14 elements are extracted and
placed at the bottom of the table. This series of
14 elements, which fits between lanthanum (La,
Z57) and hafnium (Hf, Z72) is called the
lanthanides or rare earth series. The seventh
and final period is incomplete for the moment,
but is believed to be as long as the sixth one. A
14 member series, extracted from the seventh
period and placed at the bottom of the table is
called the actinide series.
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All the atoms of the elements from group 1
possess a single outer shell electron in an s
orbital. Elements of the first group are called
alkali-metals. The atoms of the elements from
group 2 have two electrons in an outer shell.
These elements are alkaline earth metals. These
two groups are known as the s block, because
their properties arise from the presence of s
electrons. Elements of group 13 have three
electrons in their outer shell, two s electrons
and one p electron. The p electron is the
differentiating electron.
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Elements of group 14 have 4 electrons in the
outer shell (s2p2) Elements of group 15 have 5
electrons (s2p3), Elements of group 16 have 6
electrons(s2p4), Elements of group 17 have 7
electrons (s2p5) Elements of group 18 have 8
electrons(s2p6). Elements of 18 group have an
outer shell full of electrons. Because their
properties are dependent on the presence of p
electrons, groups 13 18 are called p block
elements. Elements of groups from 3 to 12 are
called the d block or transition elements and, in
a similar way, lanthanoid and actinoid elements
are f block.
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Periodic properties of the elements
The elements of the 18th group, rare gases, have
the configuration ns2 np6, except helium, whose
configuration is 1s2. That means the outer shells
of the atoms are full. These prove to be very
stable configurations and they can be altered
with great difficulty. As a result, rare gases
have a very low reactivity, they are also known
as noble gases. The electron configuration of
the elements of groups 1 and 2 differ from these
of noble gases by only one or two electrons in
the s orbital of a new shell.
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For example, the configuration of potassium and
calcium are
K Ar 4s1 Ca Ar 4s2
Except hydrogen, the elements of groups 1 and 2
are metals. The characteristic chemical
properties of metallic elements are based on the
ease of removal of one or more electrons from
their atoms to produce positive ions
K ? K e- Ca ? Ca2 2e-
Some physical properties of metals ability to
conduct heat and electricity, ductility,
malleability also arise from these distinctive
electron configurations.
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Elements of the groups 16 and 17 have an
electron configuration with two or one electron
less that those of the corresponding noble gas.
Atoms of these elements can realize the
electronic configuration of a noble gas by
gaining the appropriate number of electrons. For
example, the electron configuration of S becomes
that of Ar by gaining two electrons
S 2e- ? S2- Ne 3s23p4 Ar
The S atom becomes S2- anion. Similarly, the Cl
atom becomes Cl- anion
Cl e- ? Cl- Ne 3s23p6
Ar
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These elements whose atoms can acquire a noble
gas configuration by a small number of electrons
are non metals. Non metals are H from group
1, C from group 14, N and P from group 15, O, S
and Se from group 16 and F, Cl, Br and I from
group 17. B (13), Si, Ge (14), As, Sb (15), Tc,
Po (16) and At (17) are metalloids or semi
metals. 18th group is a special family of
elements, but noble gases may be considered non
metals. The rest of the elements, including of
course the lanthanides and actinides are metals.
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Atomic size
The sizes of atoms vary as shown in the following
figure.
Atoms get larger as we go down in a group of the
periodic table and they get smaller as we go from
left to right along a period.
Atomic radii in pm
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The atomic radius tends to increase on
descending a group due to the increment of the
number of energetic levels. The outer shell
electrons are further and further from the
nucleus, therefore less attracted by the positive
charge of the nucleus. Along a period the
charge of the nucleus increases with the atomic
number Z while the electrons are still filling
the same shell. The outer shell electrons are
attracted more strongly by the core and, as a
result, the atomic radius decreases from left to
right through a period.
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Ionic radius When electrons are removed from a
metal atom to form a positive ion (cation), a
significant reduction in size occurs. Usually,
the electrons are lost from the shell with the
highest principal quantum number. The relative
sizes of the cations are given in the figure
The relative sizes of the cations
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The ionic radius for cations in the same period
decreases from left to right. For example, in
the series of cations Na, Mg2, Al3 the number
of electrons is the same (10), while the number
of protons increases together with the atomic
number Z. Al3 is smaller that Mg2 because the
electrostatic force between the 10 electrons and
the nuclear charge of Al (13) is more powerful
than that between the 10 electrons and the
nuclear charge of Mg (12).
Na Mg2 Al3
No. of protons 11 12 13
No. of electrons 10 10 10
Ionic radius Å 0,95 0,65 0,50
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When a nonmetal atom gains one or more electrons
to form a negative ion (anion) the size
increases compared with the original atom.
The relative sizes of the anions
Along a period the ionic radius of anions
decreases from left to right.
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Ionization energy Is the energy required to
remove one electron from an individual atom in
the gaseous phase. This is the first ionization
energy
In case of metals, which have a small number of
electrons in the outer shell, a small amount of
energy is needed to remove an electron, that is
metals have low ionization energies. Inside a
group, the ionization energy tends to decrease
from top to bottom because the attraction force
of the nucleus decreases in the same way and the
electron is more easily removed.
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Nonmetals have large ionization energies because
they have a large number of electrons in the
outer shell. Nonmetals tend to gain, not to lose
electrons. Ionization energies tend to increase
from left to right along a period of the periodic
table. In general, the elements that appear in
the lower left region of the periodic table have
the lowest ionization energies and are therefore
the most chemically active metals. On the other
hand, the elements with the highest ionization
energies occur in the upper right hand region of
the periodic table. The first ionization energy
of the elements is a function of atomic number Z
85
The first ionization energy of the elements
86
The loss of the second electron occurs with
greater difficulty than the first. Therefore the
second ionization energy is higher than the first
one. A property used to describe the bond type
that results when atoms combine is
electronegativity. Electronegativity describes
the ability of an atom to attract electrons
towards itself. The most widely used
electronegativity scale was devised by Linus
Pauling. Paulings electronegativities are
dimensionless numbers ranging from about 1 for
very active metals to 4.0 for fluorine, the most
active nonmetal.
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1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
H 2,20 He
Li 0,98 Be 1,57 B 2,04 C 2,55 N 3,04 O 3,44 F 3,98 Ne
Na 0,93 Mg 1,31 Al 1,61 Si 1,90 P 2,19 S 2,58 Cl 3,16 Ar
K 0,82 Ca 1,00 Sc 1,36 Ti 1,54 V 1,63 Cr 1,66 Mn 1,55 Fe 1,83 Co 1,88 Ni 1,91 Cu 1,90 Zn 1,65 Ga 1,81 Ge 2,01 As 2,18 Se 2,55 Br 2,96 Kr 3,0
Rb 0,82 Sr 0,95 Y 1,22 Zr 1,33 Nb 1,6 Mo 2,16 Tc 1,9 Ru 2,2 Rh 2,28 Pd 2,2 Ag 1,93 Cd 1,69 In 1,78 Sn 1,96 Sb 2,05 Te 2,1 I 2,66 Xe 2,6
Cs 0,79 Ba 0,89 La 1,27 Hf 1,3 Ta 1,5 W 2,36 Re 1,9 Os 2,2 Ir 2,20 Pt 2,28 Au 2,54 Hg 2,0 Tl 1,62 Pb 2,33 Bi 2.02 Po 2,0 At 2,2 Rn
Fr 0,7 Ra 0,9 Ac 1,10
Paulings electronegativities of the elements
As a rough rule, most metals have
electronegativities of about 1.7 or less
semi-metals about 2 and nonmetals greater than 2.